The determination of activities or concentrations of ions in general and of the hydrogen ion (pH) in particular is routinely required in the chemical industry, in laboratories, and increasingly in homes (eg for the assessment of soil, pH, and swimming pool chlorination).
It has been practised to use litmus or universal indicator paper to estimate solution pH. Litmus paper is paper coated with one or more indicator dyes that change colour according to the pH of a sample in contact with the strip. The strips are then compared with a colour chart to estimate the pH. Litmus strips are relatively inexpensive and are discarded after use. The strips are easy to use, do not require calibration, and are readily portable, but have low resolution, are not useful with coloured liquids, are subject to user error, and cannot be used by colour-blind persons.
The primary method employed for accurate measurement of pH is the pH meter. The pH meter is based in principle on the measurement of electrode potential in a reversible galvanic or voltaic cell. A galvanic or voltaic cell consists of two electrodes in contact with a solution and combined in such a manner that when the electrodes are connected by an electric conductor an electric current will flow. At the interface of the electrode and the solution there exists an electrical potential call the "electrode potential". The electromotive force of the cell is then equivalent to the algebraic sum of the two electrode potentials (appropriate allowance being made for the sign of each potential difference). For reliable measurement of the EMF of such a cell it is necessary to use some form of potentiometer, for example a wheatstone bridge circuit calibrated against a reference cell of known EMF. Such cells may be chemical cells in which an overall chemical reaction takes place or "concentration" cells in which there is a change of energy due to the transfer of solute from one concentration to another. For the electrical energy produced in the cell to be related thermodynamically to the process occurring in the cell, the cell should behave reversibly in a thermodynamic sense.
Such an electrochemical cell consists of two single electrodes or half cells. For example a standard metal--metal ion electrode and a standard hydrogen electrode form a cell which is written as follows: EQU M;M.sup.+ (a=1) .function. H.sup.+ (a=1);H.sub.2( 1 atm),Pt (1)
where the vertical parallel bars indicate that the junction potential between the two different electrolytic solutions is practically eliminated by connection through a diffusion barrier, for example, a salt bridge, a frit, capillary, or the like means which allows electrical connection in solution but prevents convection and restricts diffusion of ionic species. When the concentration of the ions is such that their activity is unity, as indicated in the cell of Equation 1 by a=1, the electrode potential is designated as E.sup.0 and is equal to the standard oxidation potential. The voltage of a cell is the algebraic sum of the oxidation potential for the electrode written at the left and the reduction potential of the electrode written at the right. By convention the chemical reaction corresponding to a given cell is written so that the electrons move from left to right outside the cell.
The electrode potential changes with the activity of the ions. The fundamental equation governing the effect of activity of ions on the voltage is: ##EQU1##
where T is the temperature PA1 R=gas constant, 8.314 Joules deg.sup.-1, mole.sup.-1 PA1 F=Faraday, 96,486.7 Coulombs, equiv.sup.-1 PA1 n=number of electrons for the reaction as written PA1 Q=activity quotient ##EQU2## PA1 where E is the measured EMF for an unknown solution in the cell, R, T and F are as previously defined, pH(S) is the assigned pH value of a standard reference solution and E.sub.s is the EMF measured when this reference solution replaces an unknown solution in the cell described. PA1 a first electrode, PA1 a reference electrode, PA1 a porous membrane extending between the first electrode and the reference electrode for retaining a liquid analyte if admitted to the porous membrane in electrical contact with both said electrodes. PA1 a redox chemical system (as herein defined) contained in the porous membrane, and a diffusion barrier disposed between said electrodes.
for the generalised reaction EQU a A+b B=g G+h H (3)
It may be shown that in a single junction liquid cell with a hydrogen electrode and a reference electrode, for example a calomel electrode, the pH is defined by ##EQU3##
Three types of electrode have been used in cells of the kind under discussion for the measurement of hydrogen and other ionic concentration. The first type of reversible electrode includes a metal or a non-metal in contact with a solution of its own ions, eg zinc in zinc sulphate solution, or copper in copper sulphate solution. Non metals which, at least in principle, yield reversible electrodes are hydrogen, oxygen and the halogens, the corresponding ions being hydrogen, hydroxyl and halide ions respectively. Since the electrode materials in these cases are non-conductors, and often gaseous, finely divided platinum or other noble metal which comes rapidly into equilibrium with the hydrogen, oxygen etc has been employed for the purposes of making electrical contact. The classic hydrogen electrode which is the primary standard for all measurements of hydrogen ion concentration is an example of this first class of electrode. It consists of a small sheet of platinum coated with platinum black. In use this electrode is immersed in a solution and pure hydrogen is bubbled over the surface for at least 20-30 minutes. Although of importance as a standard, the hydrogen electrode is not practical for everyday use and is readily poisoned by species such as arsenic, heavy metals, sulphides and cyanides.
Electrodes of the second type involve a sheet or wire of metal and a sparingly soluble salt of this metal in contact with a solution of a soluble salt of the same anion. An example is an electrode consisting of silver, solid silver chloride and a solution of a soluble chloride, such as potassium chloride. These electrodes behave as if they were reversible with respect to the common anion namely the chloride ion in this case. Electrodes of this type have been made with other insoluble halides eg silver bromide and iodide and also with soluble sulphates, oxalates etc. Calomel electrodes are of this kind.
A third type of reversible electrode consists of a sheet of an inert metal, for example, gold or platinum, immersed in a solution containing both oxidised and reduced states of an oxidation reduction a system ("redox" system) for example Sn.sup.++++ and Sn.sup.++, Fe.sup.+++ and Fe.sup.++ or Fe(CN).sub.6.sup.3- and Fe(Cn).sub.6.sup.4-. The oxidised and reduced states are not necessarily ionic. For example oxidation-reduction electrodes involving organic compounds are known for example the Quinhydrone electrode. Quinhydrone is an equimolecular compound of hydroquinone, HOC.sub.6 H.sub.4 OH, and benzoquinone, OC.sub.6 H.sub.4 O. At 25.degree. C. this is dissociated to an extent of about 93% into its two components. An electrode is formed by immersing a bright inert metal such as platinum or gold in a solution (usually saturated) of Quinhydrone.
The solution needs to be freshly prepared and moreover freshly recrystallised Quinhydrone should be used in its preparation. The Quinhydrone electrode was used extensively in the 1920s but due to the inconvenience of its preparation and short solution shelf life it was rapidly superseded when the calomel electrode was developed.
The development of the glass pH electrode resulted in an electrode which far surpassed all those previously used and made practical the pH meter as now widely used. The glass electrode could be used successfully in oxidising or reducing mediums, in the presence of heavy metals, and in mixtures in which the hydrogen, quinone, and antimony electrodes would not give repetitively accurate results. In its simplest form the glass electrode consists of a thin-walled glass bulb with an inner solution in which a reference electrode is immersed. For general use the electrode is usually blown from a special glass (such as Corning 015 glass) which has a low melting point, high hygroscopicity, and relatively high electrical conductivity. The potential of a glass electrode depends upon the nature of the inner electrode and the composition of the inner solution as well as the "asymmetry potential". Because the glass electrode was suitable for mass manufacture, was relatively robust, and of long life, it became widely used for the measurement of pH. Meters utilising the glass electrode are accurate and cheap to operate. However because the glass electrode is of intrinsically high resistance an electronic voltmeter is required, and the meters are not always portable. Furthermore such pH meters require tedious calibration which involves the cost of consumable buffer solutions as well as the expenditure of valuable operator time.
PCT/GB87/00857 describes a hollow capillary fill cell having electrodes on interior surfaces of the cell and having reagents "printed" onto interior surfaces of the cell. The distance between the electrodes is significant because the cell relies on the spatial separation between the electrodes to act as a capillary or frit. That is to say the distance between the electrodes must be such that diffusion of ions from one electrode to the other takes longer than the test.